Acids, Bases & pH
Properties of Acids: Acids release a hydrogen ion into water (aqueous) solution.
1. Taste sour
 Ex: Stomach acid is hydrochloric acid. Although tasting stomach acid is
not pleasant, it has the sour taste of acid. Acetic acid is the acid
ingredient in vinegar. Citrus fruits such as lemons, grapefruit, oranges,
and limes have citric acid in the juice. Sour milk, sour cream, yogurt, and
cottage cheese have lactic acid from the fermentation of the sugar lactose.
2. Conduct an electrical current
3. Liberates H2 gas when reacted with a metal
4. Causes certain dyes to change color (Blue litmus turns red)
5. Have a pH less than 7
6. Acids corrode active metals. Even gold, the least active metal, is attacked by
an acid, a mixture of acids called 'aqua regia,' or 'royal liquid.' When an acid
reacts with a metal, it produces a compound with the cation of the metal and
the anion of the acid and hydrogen gas.
Properties of Bases: Bases release a hydroxide ion into water solution.
1. Bases taste bitter
 Ex: There are very few food materials that are alkaline, but those that
are taste bitter; ex: cough syrup, baking soda or Tums. Tasting of
bases is more dangerous than tasting acids due to the property of
stronger bases to denature protein.
2. Conduct an electrical current
3. Bases denature protein.
 This accounts for the "slippery" feeling on hands when exposed to
base. Strong bases that dissolve in water well, such as sodium or
potassium lye are very dangerous because a great amount of the
structural material of human beings is made of protein. Serious
damage to flesh can be avoided by careful use of strong bases.
4. Bases turn red litmus to blue.
5. Have a pH of more than 7
** Both Acids & Bases will loose their properties in a neutralization reaction.
Summary
Characteristics of Acids
1. Sour taste
2. Ionize when added to water. (H+)
3. Some react with metals to produce H2 gas.
4. Change the color of acid indicators.
(Litmus paper red.)
5. Neutralize bases, producing water and a salt.
6. Are electrolytes
7. Produce hydronium ions (H30+) when
dissolved in water.
1.
2.
3.
4.
5.
6.
7.
Characteristics of Bases
Bitter Taste
Slipper to touch
React with organic materials (grease & hair)
Change the color of base
indicators. (Litmus paper blue.)
Neutralize acids, producing water and a salt.
Are electrolytes
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Produce hydroxide ions (OH-) in
solution
Definitions of Acids & Bases
** There are 2 main ways to define an acid & a base. In chemistry, the proper
definition of an acid & base is taken from the combination of each definition.
Arrhenius Definitions: (Traditional)
Acid: substance whose aqueous solution contains H+ ions.
Ex: HCl, HNO3, H2SO4
Base: substance whose aqueous solution contains OH- ions.
Ex: NaOH, Ca(OH)2
Examples: HCl
HNO3
H2SO4
Strong Acid: HCl(g)
[H3O+ hydronium]
Weak Acid: HC2H3O2 (l)
Examples: NaOH
Ca(OH)2
Strong Base: NaOH
Weak Base: HC2H3O2
2
OR
ACIDS
BASES
H2O
HCl
H2O
H+(aq) + Cl-(aq)
Na+(aq)
NaOH
+
OH-(aq)
Neutralization: As you can see above, acids release H+ into solution and bases
release OH-. If we were to mix equal amounts of a strong acid and base together,
the H+ ion would combine with the OH- ion to make the molecule H2O, or plain
water:
H+(aq) +
OH-(aq)
H2O
The neutralization reaction of an acid with a base will always produce water and a
salt, as shown below:
Acid
Base
HCl
+ NaOH
H2O
+ NaCl
HBr
+
H2O
+
KOH
Water
Salt
KBr
(Salt: Generally, any ionic compound except those that contain hydroxide or
hydrogen ions. Specifically, any compound other than water formed by the
reaction of an acid and a base.)
While Arrhenius helped explain the fundamentals of acid/base chemistry,
unfortunately his theories have limits. For example, the Arrhenius definition does
not explain why some substances like common baking soda (NaHCO3) or
ammonia (NH3) can act like a base even though they do not contain hydroxide
ions.
Acidic Solutions
Basic Solutions
Neutral Solutions
[H+] > [OH-]
[H+] < [OH-]
[H+] = [OH-]
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In 1923, the Danish scientist Johannes Brønsted and the Englishman Thomas
Lowry published independent yet similar papers that refined Arrhenius' theory.
Brønsted – Lowry Definition:
Acid: substance that can donate a proton (H+)
Base: substance that can accept a proton (H+)
Ex:
Acid
HCl
Base
+ NaHCO3
Salt
⇌
H2CO3 +
NaCl
⇌is used as an equilibrium sign.
The rxn does not go to
completion; Reactants partially dissociate (break apart). How
much/how quickly they dissociate depends on if they are weak or
strong! Key: the stronger an acid/base is, the more it dissociates.
Brønsted – Lowry Acid-Base Pairs: A conjugate pair is an acid-base pair that
differs by one proton in their formulas (remember: proton, hydrogen ion, etc.).
Conjugate Acid: what the base use to be plus a H+
Conjugate Base: what the acid use to be without the H+
Ex: HCl + NH3 ⇌ NH4 + ClAcid + Base ⇌ CA + CB
Ex: H2CO3 + CO2+ ⇌ HCO2+ + HCO3Acid + Base ⇌ CA
+ CB
Ex: HNO3 + H2O ⇌ H3O+ + NO3¯
The acids are HNO3 and H3O+ and the bases are H2O and NO3¯.
Another way to look at an acid-base reaction is to think of it as a competition
between two bases (think about it!) for a proton. If the stronger of the two acids
and the stronger of the two bases are reactants (appear on the left side of the
equation), the reaction is said to proceed to a large extent.
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If the weaker of the two acids and the weaker of the two bases are reactants (appear
on the left side of the equation), the reaction is said to proceed to only a small
extent
Self Ionization of Water: the reaction in which two water molecules react to give
ions; also called the dissociation of water.
Kw = equilibrium constant of pure water = 1.0 x 10-14 at 250C (Standard Temp)
Dissociation of Water Web Site: H2O + H2O ⇌ H3O+ + OH Amphoteric- substance whose solution can have both acidic & basic properties
Ex: H2O, HSO4-, HPO4-2, H2PO4 Ex: HNO3 + H2O
Ex: NH3 + H2O
Strong acids and bases ionize (separate into ions) completely. While weak A/B
ionize partially
H2O
 Ex: Strong acid ionizing
HCl
H+(aq) + Cl-(aq)
H2O
 Ex: Weak acid ionizing
HC2H3O2
⇌
H+(aq) + C2H3O2-(aq)
Strength vs Concentration
Strength is how fully the acid or base
will ionize. A compound that ionizes
90% is stronger than a compound that
ionizes only 20%.
Common Strong Acids/Bases You
Need to Memorize! (The BIG 6)
Acids
Bases
HF
LiOH
HCl
NaOH
HBr
KOH
HI
Mg(OH)2
H2SO4
Ca(OH)2
HNO3
Ba(OH)2
Concentration tells us the amount of acid
or base in a solution. In a “concentrated”
acid solution, there is lots of acid
compared to water. In a “dilute” acid
solution, the acid is watered down. In
other words, there is more water than
acid. Concentration is expressed in terms
of Molarity.
Molarity = moles / liters
** The categories strong/weak still contain a wide range of strengths
** Weak is not insignificant – most reactions within living systems involve weak
acids or bases
** Strength is not concentration – can have a diluted strong acid/base or
concentrated weak acid/base
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Strength vs. Concentration
Strength: How easy an Acid or Base will break apart
into its ions. Things that won’t break apart into ions
are molecular
How can we test to see if a substance should be
classified as strong or weak?
Qualitatively…use electricity.
Stronger Acids or Bases have more ions in solution
so they conduct electricity better.
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Quantitatively…use Equilibrium Constant, K
**(All K’s, Equilibrium Constants, are temp dependent. All discussions
regarding them are assumed to be at STP, 250C, unless otherwise noted)
Definitions:
→ a reaction that is complete (100% ionization)
⇌ a reaction that only takes place partially (< 100%
ionization)
[ ] the concentration of the compound (measured using
molarity)
Keq the equilibrium constant for an reaction at equilibrium
For aA + bB ⇌ cC + dD
[products] x C D


y
[reactants ]
Aa Bb
c
K eq
d
*** Keq > 1, products are favored at equilibrium
*** Keq < 1, reactants are favored at equilibrium.
Ex. BrCl(g) ⇌
Cl2(g) + Br2(g)
Keq =
7
Keq ex1 : N2O4 and NO2 exist in equilibrium according to
the following equation: N2O4(g) ⇌ NO2(g)
A liter of a gas mixture at 100C at equilibrium contains
0.0045 moles of dinitrogen tetroxide and 0.030 moles of
nitrogen dioxide.
a. BCE:?
b. Write the expression for the equilibrium constant
c. Calculate the equilibrium constant for the rxn.
Keq ex2 : Bromine chloride decomposes to form bromine and
chlorine. At a certain temp, the Keq for the reaction is 11.1, and
the equilibrium mixture contains 4.00 moles Cl2. How many
moles of Br2 and BrCl are present in the equilibrium mixture?
a. BCE:
b. Keq = ?
c. Plug and chug!
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Keq ex3 (trickie) :One mol of hydrogen gas and 1.00 mole of
iodine are sealed in a 1-L flask and allowed to react at 4500C.
At equilibrium 1.56 mole of hydrogen iodide is present.
Calculate the Keq for the rxn.
a. BCE?:
b. Keq = ?
c. Plug and chug! (Hint: calculate how much of each
reactant remain at equilibrium.)
Ka the acid dissociation constant; used when acids are placed
in water
Ka 
[dissociat ed acid]
[associate d acid]
* Generic Ka Equation: HA + H2O ⇌ H3O+ + AOR
HA + H2O ⇌ H+ + AKa =
[H3O+] [A-]
[HA]
 Ka = [H+] only for strong acids! (WA’s – solve for Ka)
 Note that water is omitted from the expression because it is present in such
vast excess that its concentration changes negligibly on the formation of
equilibrium and is therefore effectively constant. The concentration of the
water is included in the equilibrium constant, and Ka can be thought of as a
modified equilibrium constant.
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 The greater the Ka, the greater the dissociation and thus the lower the pH or
more acidic a solution is!
 Ex: The Ka for HCl is 100,000,000 whereas the Ka for acetic acid is
0.000018
Kb the base dissociation constant; used when bases are placed
in water
[conjugate acid][OH - ]
Kb 
[base]
* Generic Kb Equation: NH3 + H2O ⇌ NH4+ + OH[NH4+] [OH-]
Kb =
[NH3]
 Kb = [OH-] only for strong bases (WB’s – solve for Kb)
 The greater the Kb, the greater the dissociation and thus the higher the pH
or more basic a solution is!
3 Steps for solving for equilibrium concentrations and K’s
1. Write a general equation.
2. Setup an ICE Table (Initial concentration, Change in
concentration, Equilibrium concentration)
3. Write a K expression and plug in the equilibrium
concentrations.
Ex. Aspirin is a weak acid. A chemist mixes 0.1000 moles of
aspirin with water to make 1 L of solution. At equilibrium the
H3O+ concentration was found to be 0.0057 M. What is the Ka
for aspirin?
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Ex: Do questions 1 & 4 from Special K’s worksheet, after
viewing the following websites:
Correct annunciations: http://antoine.frostburg.edu/chem/senese/101/acidbase/glossary.shtml
Disregarding the effects of [H3O+] and [A-]:
http://www.shodor.org/UNChem/basic/ab/index.html
http://www.chem.tamu.edu/class/majors/tutorialnotefiles/pka.htm
Kw ion product constant for water
Kw = [H+][OH-] = 1.0 x 10-14 (mol/L)2
Kw = Ka x Kb
The pH Scale
 Developed in 1909 by Soren Sorenson
 Developed to classify acidic & basic solutions
 pH: mathematical scale in which the concentration of hydronium ions in a
solution is expressed as a number ranging from 0 to 14; 0 acidic & 14 basic
 Based on the number 10; a pH of 3 is 10x as acidic as a pH of 4 & a pH of 3 is
____x as acidic as a pH of 5.
 The more hydronium ions present, the more acidic the solution
 The true math scale ranges from 100 to 10-14
[H+]
1 x 100
1 x 10-1
1 x 10-7
1 x 10-11
1 x 10-13
1 x 10-14
[OH-]
1 x 10-14
1 x 10-13
1 x 10-7
1 x 10-3
1 x 10-1
1 x 100
pH
0
1
7
11
13
14
pOH
14
13
7
3
1
0
substance
1M HCl
0.1M HCl
Pure water
Household ammonia
0.1 M NaOH
1M NaOH
Ex: a solution with a pH of 9 has 10-9 hydronium ions present in solution whereas
a solution with a pH of 3 has 10-3 hydronuium ions present in solution
 Ex: a solution whose hydrogen ion concentration is 0.001M, or 1x10-3, you
would simply say the solution has a pH of 3.
 Uses of pH
 Pool chemicals
 Soil
 Rivers, lakes, streams,…
 Personal care products
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 Cleaners
 Tooth decay by the pH of saliva changing due to bacteria adjusting pH levels
 Formation of Acid Rain (a series of rxn’s must occur)
Ex: SO2(g) + H2O(l) → H2SO3(aq)
SO3(g) + H2O(l) → H2SO4(aq) Sulfuric Acid
 The pOH scale is based on the pH scale & uses the fact that the sum of
hydronium ions + hydroxide ions = 14; in other words, pOH = 14 – pH
 Can you make a pOH scale?
Calculating pH
 pH is a convenient way to measure the acidity [H+] or causticity [OH-]of a
solution.
 pH measures the amount of H+ ions present in a solution.
More H+ ions 
pH
+
Less H ions 
pH
 pOH measures the amount of OH ions present in a solution
Summary of equations to use
pH = -log [H+]
[H+] =10-pH
pOH = -log [OH-]
[OH-]=10-pOH
pH + pOH = 14
[H+][OH-] = 1 x 10-14
Ex. 1
What is the pH of a solution whose [H+] = 2.50 x 10-12 M?
Ex.2
What is the pH of a solution whose [H+] = .0015 M?
Ex. 3
pH paper measures a substance to have a pH = 1.2. What is the [H+]?
Ex 4
pH of a solution is 3.45. Calculate the pOH, [OH-], and [H+].
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